Transcript Chapter 4

Chapter 5
The Periodic Law
Section 5-1
History of the Periodic Table
Stanislao Cannizzaro
13 Jul 1826 – 10 May 1910
Cannizzaro
• Was the first scientist to accurately
measure atomic masses.
• This was incredibly important for the work
that Dmitri Mendeleev was going to do.
Dmitri Mendeleev
8 Feb 1834 – 2 Feb 1907
Dmitri Mendeleev
• Arranged the elements on cards in order
of increasing atomic mass
• Found columns of elements with similar
properties
• There were gaps in his columns
Basic Version of Mendeleev’s
Periodic Table
The gaps
• Hypothesized the gaps were undiscovered
elements.
• Predicted the props of these elements.
• Predicted them well.
Mendeleev’s Mistake
• There were irregularities when arranged
according to atomic weight.
Henry Mosely
23 Nov 1887 – 10 Aug 1915
Henry Mosely
• Discovered a unique charge on the
nucleus of the atom.
• Arranged the elements according to
increasing atomic number
• When he did this, the irregularities
disappeared.
Periodic Law
• The properties of elements tend to change
with atomic number gradually, in a
periodic way.
John Strutt
12 Nov 1842 – 30 Jun 1919
William Ramsay
2 Oct 1852 – 23 Jul 1916
Strutt & Ramsay
• In 1894, they discovered argon.
• Nobody noticed it before because it is
completely unreactive.
• In 1868, helium had been discovered as
part of the sun and in 1895, Ramsay
showed its existence on earth.
• In 1898, Ramsay discovered krypton and
xenon. Friedrich Ernst Dorn discovered
radon in 1900.
Lanthanides
• Discovered in the early 1900’s.
• They are found in the f block
• They are shiny and act like the alkaline
earth metals.
Actinides
• All of these elements are radioactive.
• They are found in the f block
Alkali Metals
• Group 1
• These elements are soft and can be cut
with a knife.
• They are highly reactive. The will react
with both air and water.
• They form alkaline/basic solutions (the
opposite of acidic solutions).
• Their electron configurations all end s1.
Sodium
Alkaline Earth Metals
• Group 2 on the periodic table.
• These elements are harder and denser
than the alkali metals.
• They are also reactive, but less so than
the alkali metals.
• They will also form alkaline/basic
solutions.
• Their electron configurations all end s2.
Magnesium
Hydrogen
• This element doesn’t belong with any
group.
• Its electron configuration is 1s1.
Helium
• Even though its electron configuration
ends s2, it isn’t an alkaline earth metal. It
is a noble gas because its highest energy
level orbitals are full.
Transition metals
• AKA transition elements
• AKA d block elements
• These elements are what we typically
picture as common metals.
• Their d orbitals are being filled.
Transition metals cont’d
• They are shiny and good conductors of
electricity.
• They are less reactive than the other
metals.
• Some like gold are highly unreactive.
Chromium metal
Main Group Elements
• Properties of these elements vary greatly
because they include metals, nonmetals,
metalloids, and noble gases.
• They include the elements of the s and p
blocks.
• Sometimes they are called the
representative elements because metals,
nonmetals, metalloids, and noble gases
are all represented
Halogens
• These are the most reactive nonmetals.
• They have 7 electrons in the outermost
energy level and their electron
configurations all end in s2p5.
• They will react with metals to form salts.
Noble Gases
•
•
•
•
Group 18
No stable compounds for He, Ne, or Ar
Very low reactivity for the rest
Full s & p orbitals (s2p6) in the higest
energy level
– This is very stable - they have no need to
react with anything else.
Noble Gases cont’d
• Most other atoms gain/lose e- to achieve
this e- configuration
• Ne & Ar are used in signs
• He - Low density - Air ships & weather
balloons
Noble Gases Cont’d
• Rn - Radioactive
• Found in homes - Linked to Lung CA
• Once you test for it, you must disclose the
results to potential buyers
5 – 3 Electron Configuration and
Periodic Properties
Atomic Radius
• As you move down a group it increases
• The outermost e- are being added to
higher energy levels (further from the
nucleus.
Atomic Radius
• As you move across a period, it decreases
• Even though e- are being added, they are
added to the same energy level (same
distance from the nucleus).
Atomic Radius
• The charge on the nucleus increases as
you move across the period and so it has
a “tighter” hold on the e- being added.
Shielding Effect
• The reduction of the attractive force
between a nucleus and its outer electrons
due to the blocking effect of inner
electrons.
Ionization Energy
• The amount of energy needed to remove
an electron from an atom
Ion
• An atom that has gained or lost an e• If it has gained an e-, it will be _____.
• If it has lost an e-, it will be _____.
Ionization Energy
• As you go down a group, it decreases
• Shielding effect and electrons are being
added to higher energy levels.
Ionization Energy
• As you move across a period, it increases
• The charge on the nucleus increases as
you move across the period and so it has
a “tighter” hold on the e- being added.
Electron Affinity
• Measures the tendency of an atom to attract
electron
• The energy change that occurs when an
electron is acquired by a neutral atom.
• Metals tend to have positive energy changes,
they do not have a tendency to attract electrons
• Non-metals tend to have negative energy
changes, they have a strong tendency to attract
electrons
EA
• As you move down a group, electrons add
with greater difficulty (values become
more positive)
• Shielding Effect and electrons are being
added to higher energy levels.
• There are exceptions
EA
• Electrons add more easily as you move
across the periodic table.(values become
more neg)
• The charge on the nucleus increases as
you move across the period and so it has
a “tighter” hold on the e- being added.
Electronegativity
• Tendency for an atom to attract e- to itself
when combined with another atom.
• F is the most EN
• EN decreases as you move down a group
• EN increases as you move across a
period
Electronegativity
• Based on the Pauling
Scale.
• Linus Pauling
• 28 Feb 1901 –
19 Aug 1994